CLASSICAL any amount of energy can be emitted or absorbed NON-CLASSICAL energy can be emitted or absorbed only in discrete quantities (little packages) energy is not continuous QUANTUM smallest amount of energy which can be absorbed/ emitted
Albert Einstein Electromagnetic radiation can be viewed as a stream of particle-like units called photons POSTULATE 3. Properties of Photons
E initial E final ABSORPTION OF A PHOTON atoms and molecules absorb discrete photons (light quanta)
de Broglie Duality of Wave and Corpuscle light has properties of a wave and of a particle
SUMMARY 1. light can be described as a wave of a wavelength and frequency 2. light can be emitted or absorbed only in discrete quantities (quantum - photon) 3. duality of wave and corpuscle
de Broglie wavelength each particle can be described as a wave with a wavelength λ 4. Properties of Electrons
matter and light (photons) show particle and wave-like properties WAVE-PARTICLE DUALITY MASS INCREASESWAVELENGTH GETS SHORTER MASS DECREASES WAVELENGTH GETS LONGER
Wave-likeParticle-like BaseballProtonPhotonElectron WAVE-PARTICLE DUALITY large pieces of matter are mainly particle-like small pieces of matter are mainly wave-like MASS
1. light behaves like wave and particle 2. electron behaves like wave and particle 3. electrons are constituents of atoms 4. light is emitted/absorbed from atoms in discrete quantities (quanta)
E initial E final EMISSION OF A PHOTON atoms and molecules emit discrete photons electrons in atoms and molecules have discrete energies 5. Electrons, Photons, Atoms
EMISSION SPECTRA analyze the wavelengths of the light emitted only certain wavelengths observed only certain energies are allowed in the hydrogen atom
Balmer found that these lines have frequencies related n = 1, 2, 3, 4, 5…
electrons move around the nucleus in only certain allowed circular orbits e-e- THE BOHR ATOM each orbit has a quantum number associated with it QUANTUM NUMBERS n is a QUANTUM NUMBER n= 1,2,3,4……... n = 4 n = 3 n = 2 n = 1
n = 4 n = 3 n = 2 n = 1 THE BOHR ATOM QUANTUM NUMBERS and the ENERGY Z = atomic number of atom A = 2.178 x 10 -18 J = Ry THIS ONLY APPLIES TO ONE ELECTRON ATOMS OR IONS
BOHR ATOM ENERGY LEVEL DIAGRAM Z=1 HYDROGEN ATOM!
EnEn ENERGY n=1 -A BOHR ATOM ENERGY LEVEL DIAGRAM
n=1 -A n=2 -A/4 EnEn ENERGY BOHR ATOM ENERGY LEVEL DIAGRAM
n=1 -A n=2 -A/4 EnEn n=3 -A/9 n=4 ENERGY BOHR ATOM ENERGY LEVEL DIAGRAM e-e- E photon = h ELECTRON EXCITATION
n=1 -A n=2 -A/4 EnEn 0 n=3 -A/9 n=4 Energy e-e- ELECTRON DE-EXCITATION emission of energy as a photon e-e-
nini nfnf only a photon of the correct energy will do ABSORPTION OF A PHOTON
nfnf nini EMISSION OF A PHOTON This means energy is emitted!
nini nfnf IONIZATION OF AN ATOM This means energy is absorbed!
EE the ionization energy for one mole is IONIZATION ENERGY = 2.178x 10 -18 J atom -1 x 6.022x10 23 atoms mol -1 =13.12 x 10 5 J mol -1 = 1312 kJ mol -1 = 2.178 x 10 -18 J for one atom
e-e- THE BOHR ATOM QUANTUM NUMBERS n = 4 n = 3 n = 2 n = 1 absorption emission ionization energy
6. HEISENBERG’S UNCERTAINTY PRINCIPLE x is the uncertainty in the particle’s position p is the uncertainty in the particle’s momentum in the microscopic world you cannot determine the momentum (velocity) and location of a particle simultaneously EXP5
THE HEISENBERG UNCERTAINTY PRINCIPLE if particle is big thenuncertainty small EXP5
This means we have no idea of the velocity of an electron if we try to tie it down! Alternatively if we pin down velocity we have no idea where the electron is! So for electrons we cannot know precisely where they are!
we cannot describe the electron as following a known path such as a circular orbit Bohr’s model is therefore fundamentally incorrect in its description of how the electron behaves. we cannot know precisely where electrons are!
orbit of an electron at radius r (Bohr) probability of finding an electron at a radius r (Schroedinger, Born)
1. Schroedinger defines energy states an electron can occupy 2. square of wave function defines distribution of electrons around the nucleus high electron density - high probability of finding an electron at this location low electron density - low probability of finding an electron at this location atomic orbital wave function of an electron in an atom each wave function corresponds to defined energy of electron an orbital can be filled up with two electrons (box) EXPVII
most atoms have more than two electrons each electron in an atom is different electrons have different ‘labels’ called quantum numbers
QUANTUM NUMBERS 1.principle quantum number 2. angular momentum quantum number 3. magnetic quantum number 4. spin quantum number
1. principle quantum number n n = 1, 2, 3, 4, 5… hydrogen atom: n determines the energy of an atomic orbital measure of the average distance of an electron from nucleus n increases → energy increases n increases → average distance increases
e-e- n = 4 n = 3 n = 2 n = 1 n = 1 2 3 4 5 6 K L M N O P ‘shell’ maximum numbers of electrons in each shell 2 n 2 EXP6
2. angular momentum quantum number l = 0, 1, … (n-1) l = 0 1 2 3 4 5 s p d f g h define the ‘shape’ of the orbital
ORBITALS AND QUANTUM NUMBERS 1.principle quantum number 2. angular momentum quantum number 3. magnetic quantum number 4. spin quantum number n = 1, 2, 3, 4, 5… l = 0, 1, … (n-1) m l = -l, (-l + 1), … 0…… (+l-1) +l m s = -1/2; + 1/2
(n, l, m l, m s ) ATOMIC ORBITALS nlmlml orbitalsdesignation 10011s 20012s 1-1,0,+132p x,2p y,2p z 30013s 1-1,0,+133p x,3p y,3p z 2-2,-1,0,+1,+253d xy,3d yz,3d xz, 3d x 2 -y 2,3d z 2 4…………
H Atom Orbital Energies energy level diagram H atom 3s3p3d2s2p1s3s3p3d2s2p1s E energy depends only on principal quantum number orbitals with same n but different l are degenerate
1s1s E 2s2s 2p2p 3s3s 3p3p 3d3d 4s4s 4p4p 5s5s 4d4d MULTI-ELECTRON ATOM orbitals with same n and different l are not degenerate energy depends on n and m l EXAMPLES [Xe]
Periodic Table of the Elements period groupgroup chemical reactivity - valence electrons ns 1 ns 2 ns 2 np 6 ns 2 (n-1)d x
cations are smaller than their atoms anions are larger than their atoms Na is 186 pm and Na + is 95 pm F is 64 pm and F - is 133 pm same nuclear charge and repulsion among electrons increases radius one less electron electrons pulled in by nuclear charge O < O – < O 2–